Changes of State & Internal Energy
Changes of State
Matter can change from one state to another when it is heated or cooled. These changes of state are physical changes, not chemical changes. The material keeps its original properties and the change can be reversed. For example, water can freeze into ice and then melt back into water — the substance is still water throughout.
The key changes of state are:
Melting — solid to liquid. Freezing — liquid to solid. Evaporation — liquid to gas (from the surface, at any temperature). Boiling — liquid to gas (throughout the liquid, at a specific temperature). Condensation — gas to liquid. Sublimation — solid directly to gas (without passing through the liquid state).
During any change of state, mass is conserved. No atoms are created or destroyed — the same particles are simply rearranged into a different structure. If you freeze 200 g of water, you get exactly 200 g of ice.
Changes of state are physical changes — they are reversible and mass is conserved. The six changes are: melting, freezing, evaporation, boiling, condensation, and sublimation.
Internal Energy and Heating
Every system stores energy in its particles. This internal energy is the total kinetic energy (from particle movement) and potential energy (from the forces between particles) of all the particles in the system.
When you heat a substance, you increase its internal energy. This can have two effects:
Temperature rises: If the substance is not changing state, the added energy increases the kinetic energy of the particles — they move faster (in liquids and gases) or vibrate more vigorously (in solids). The temperature goes up.
Change of state: If the substance is at its melting or boiling point, the added energy goes into breaking or weakening the bonds between particles (increasing potential energy), not into making them move faster. The temperature stays constant during the change of state, even though energy is still being transferred in.
Temperature is a measure of the average kinetic energy of the particles. During a change of state, all the energy being supplied goes into overcoming the forces of attraction between particles (increasing potential energy), not into increasing their speed. Since kinetic energy does not increase, the temperature remains constant until the change of state is complete.
Specific Heat Capacity
Specific heat capacity tells you how much energy is needed to raise the temperature of 1 kg of a substance by 1 \(^{\circ}\)C. Different materials have different specific heat capacities. Water has a high specific heat capacity (4200 J/kg \(^{\circ}\)C), meaning it takes a lot of energy to heat up — but it also releases a lot of energy when it cools down, which is why water is used in central heating systems.
The equation is:
$$\Delta Q = m \times c \times \Delta \theta$$
where \(\Delta Q\) is the change in thermal energy in joules (J), \(m\) is mass in kilograms (kg), \(c\) is specific heat capacity in J/kg \(^{\circ}\)C, and \(\Delta \theta\) is the change in temperature in \(^{\circ}\)C.
For example, to heat 2 kg of water from 20 \(^{\circ}\)C to 100 \(^{\circ}\)C: \(\Delta Q = 2 \times 4200 \times (100 - 20) = 2 \times 4200 \times 80 = 672{,}000\) J (or 672 kJ).
\(\Delta Q = m \times c \times \Delta \theta\). Specific heat capacity (\(c\)) is the energy needed to raise the temperature of 1 kg of a substance by 1 \(^{\circ}\)C. Water: \(c\) = 4200 J/kg \(^{\circ}\)C.
Specific Latent Heat
Specific latent heat is the energy required to change the state of 1 kg of a substance without changing its temperature. There are two types:
Specific latent heat of fusion — the energy needed to change 1 kg of a substance from solid to liquid (or liquid to solid) at its melting point.
Specific latent heat of vaporisation — the energy needed to change 1 kg of a substance from liquid to gas (or gas to liquid) at its boiling point. The latent heat of vaporisation is always larger than the latent heat of fusion for the same substance, because the particles must completely overcome all forces of attraction to become a gas.
The equation is:
$$Q = m \times L$$
where \(Q\) is thermal energy for the change of state in joules (J), \(m\) is mass in kilograms (kg), and \(L\) is specific latent heat in joules per kilogram (J/kg).
To determine the specific heat capacity of water: place a known mass of water in an insulated container with an immersion heater. Record the starting temperature, switch on the heater for a measured time, and record the final temperature. Use the power rating of the heater and the time to calculate the energy supplied (\(E = P \times t\)), then rearrange \(\Delta Q = m \times c \times \Delta \theta\) to find \(c\).
For the temperature–time graph of melting ice: place crushed ice in a beaker and heat it steadily, recording the temperature every 30 seconds. The graph shows temperature rising, then a flat section at 0 \(^{\circ}\)C (while ice melts), then rising again. The flat section demonstrates that energy is being used to break bonds (latent heat) rather than raising temperature.
Reducing Unwanted Energy Transfer
Thermal insulation reduces unwanted energy transfer from warm objects to their surroundings. Houses lose thermal energy through walls, roofs, windows, and floors. Insulation methods include cavity wall insulation (trapping air in foam to reduce convection and conduction), loft insulation (fibreglass traps pockets of air), double glazing (two panes with a gap reduce conduction and convection), and draught excluders (reduce convection through gaps).
The principle behind all these methods is the same: they trap air. Air is a poor conductor, and when trapped in small pockets it cannot circulate (convect), making it an excellent insulator.
Gas Pressure and Temperature
The pressure of a gas is caused by particles colliding with the walls of their container. Each collision exerts a tiny force on the wall. The total of all these forces, divided by the area of the wall, gives the pressure.
When the temperature of a gas increases (at constant volume), the particles gain kinetic energy and move faster. They hit the walls more frequently and with greater force, so the pressure increases. Conversely, cooling a gas reduces particle speed and lowers the pressure.
Gas pressure is caused by particles colliding with container walls. Heating a gas at constant volume makes particles move faster, collide more frequently and with more force, so pressure increases.
Absolute Zero and the Kelvin Scale
Absolute zero is the lowest possible temperature: −273 \(^{\circ}\)C (or 0 K on the Kelvin scale). At absolute zero, particles have no kinetic energy and stop moving entirely. It is impossible to cool anything below this temperature.
To convert between Celsius and Kelvin:
$$\text{Temperature in K} = \text{Temperature in } ^{\circ}\text{C} + 273$$
For example, 100 \(^{\circ}\)C = 373 K, and 0 K = −273 \(^{\circ}\)C.
The Kelvin scale is an absolute scale that starts at absolute zero, where particle motion stops. This makes it directly proportional to the kinetic energy of particles. If you double the temperature in kelvin, you double the average kinetic energy. This relationship does not work with Celsius because 0 \(^{\circ}\)C is not the point of zero kinetic energy.
When explaining why temperature stays constant during a change of state, always link it to potential energy increasing (bonds breaking or forming) while kinetic energy stays the same. Simply saying “energy is used to change state” is too vague.
Summary
Changes of state are reversible physical changes where mass is conserved. Heating can raise temperature (increasing kinetic energy) or cause a change of state (increasing potential energy at constant temperature). Specific heat capacity (\(\Delta Q = mc\Delta\theta\)) measures energy to change temperature. Specific latent heat (\(Q = mL\)) measures energy to change state. Gas pressure depends on particle collisions, and absolute zero (−273 \(^{\circ}\)C) is the point where particle motion stops.